pH Calculator
Calculate the pH of any solution — strong acids, strong bases, weak acids, or buffer solutions. Also converts between pH, pOH, [H⁺], and [OH⁻].
How to Calculate pH
The method depends on the type of solution:
- From hydrogen ion concentration: pH = -log₁₀[H⁺]. Example: [H⁺] = 0.001 M → pH = -log(0.001) = 3.
- Strong acid (fully dissociates): [H⁺] = molarity × n, where n is the number of H⁺ ions released. Example: 0.1 M HCl → [H⁺] = 0.1 → pH = 1.
- Weak acid: pH = ½(pKa - log₁₀ C). Example: 0.1 M acetic acid (pKa = 4.74) → pH = ½(4.74 + 1) = 2.87.
pH Formulas
| Calculation | Formula |
|---|---|
| pH from [H⁺] | pH = -log₁₀[H⁺] |
| [H⁺] from pH | [H⁺] = 10-pH |
| pOH from pH | pOH = 14 - pH |
| pH of strong acid | pH = -log₁₀(M × n) |
| pH of strong base | pH = 14 + log₁₀(M × n) |
| pH of weak acid | pH = ½(pKa - log₁₀ C) |
| pH of buffer | pH = pKa + log₁₀([A⁻]/[HA]) |
Where M = molarity, n = number of ionizable H⁺ or OH⁻, C = acid/base concentration, Ka = acid dissociation constant, pKa = -log₁₀(Ka).
The pH Scale
The pH scale runs from 0 (most acidic) to 14 (most basic), with 7 as neutral. Each whole number is a tenfold change in hydrogen ion concentration — pH 3 is ten times more acidic than pH 4.
| pH | Category | Example |
|---|---|---|
| ~0 | Strongly acidic | Battery acid (1 M H₂SO₄) |
| ~2 | Acidic | Lemon juice, gastric acid |
| ~3 | Acidic | Vinegar, orange juice |
| ~5 | Weakly acidic | Coffee, acid rain |
| ~6 | Weakly acidic | Urine, saliva, milk |
| 7 | Neutral | Pure water (25°C) |
| ~8 | Weakly basic | Sea water, blood |
| ~9 | Basic | Baking soda solution |
| ~11 | Basic | Ammonia, soapy water |
| ~13 | Strongly basic | Bleach, drain cleaner |
| ~14 | Strongly basic | 1 M NaOH |
Common Ka and pKa Values
Reference values for common weak acids at 25°C. Use these with the weak-acid mode above.
| Acid | Formula | Ka | pKa |
|---|---|---|---|
| Acetic acid | CH₃COOH | 1.8 × 10⁻⁵ | 4.74 |
| Hydrofluoric acid | HF | 6.8 × 10⁻⁴ | 3.17 |
| Hydrocyanic acid | HCN | 4.0 × 10⁻¹⁰ | 9.40 |
| Carbonic acid (Ka₁) | H₂CO₃ | 4.3 × 10⁻⁷ | 6.37 |
| Formic acid | HCOOH | 1.77 × 10⁻⁴ | 3.75 |
| Hypochlorous acid | HOCl | 6.2 × 10⁻⁸ | 7.21 |
| Nitrous acid | HNO₂ | 4.5 × 10⁻⁴ | 3.35 |
| Sulfurous acid (Ka₁) | H₂SO₃ | 1.1 × 10⁻² | 1.96 |
| Phosphoric acid (Ka₁) | H₃PO₄ | 7.5 × 10⁻³ | 2.12 |
| Benzoic acid | C₆H₅COOH | 6.3 × 10⁻⁵ | 4.20 |
Frequently Asked Questions
What is pH?
pH measures hydrogen ion concentration on a logarithmic scale from 0 to 14. pH 7 is neutral, below 7 is acidic, above 7 is basic. The formula is:
pH = -log₁₀[H⁺]
where [H⁺] is the hydrogen ion concentration in mol/L.
How do you calculate pH?
Use pH = -log₁₀[H⁺]. For strong acids with known molarity: [H⁺] = molarity × number of H⁺ ions released.
Example: For 0.1 M HCl, [H⁺] = 0.1 M, so pH = -log(0.1) = 1.
What is the difference between pH and pOH?
pH measures hydrogen ion concentration [H⁺]; pOH measures hydroxide ion concentration [OH⁻]. They are complementary: at 25°C,
pH + pOH = 14
for any aqueous solution.
How do you calculate the pH of a weak acid?
Use the approximation
pH = ½ (pKa - log₁₀ C)
where pKa = -log₁₀(Ka) and C is the acid concentration. This assumes Ka << C.
Example: 0.1 M acetic acid (Ka = 1.8 × 10⁻⁵, pKa = 4.74) → pH = ½(4.74 - log 0.1) = ½(4.74 + 1) = 2.87.
What is the Henderson-Hasselbalch equation?
The Henderson-Hasselbalch equation calculates the pH of a buffer solution:
pH = pKa + log₁₀ ([A⁻] / [HA])
where [A⁻] is the conjugate base concentration and [HA] is the weak acid concentration. A buffer has maximum capacity when [A⁻] = [HA], giving pH = pKa.
Can pH be negative or above 14?
Yes, technically. Very concentrated strong acids (above 1 M) can have pH below 0, and very concentrated strong bases can have pH above 14. The conventional 0–14 range applies to dilute aqueous solutions at 25°C.
Related Chemistry Tools
- Molarity Calculator — find the concentration needed for a pH calculation.
- Dilution Calculator (M₁V₁ = M₂V₂)
- Molar Mass Calculator
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All calculations on CoolConversion use formulas and constants documented by internationally recognised standards bodies (such as IUPAC and NIST).
